Hydrates

Hydrates usually have stoichiometric amounts of water. This means that instead of just containing tiny amounts of water, they contain an amount large enough to be comparable to the amount of other elements present in the compound.

The general way we write hydrates is:

$$AB\circ nH_2O$$

Where AB is the formula for the non-water (anhydrous) part of the complex, and n is the number of moles of water.

For example, here is magnesium sulfate heptahydrate:

$$MgSO_4\circ 7H_2O$$

Here we see that for every one mole of magnesium sulfate (MgSO₄), there are 7 moles of water.

Hydrates Formation

Water is all around, even in the air as water vapor. When some chemicals are exposed to the air, they can either absorb it onto their surface (like a sponge) or incorporate it into their structure.

So, how is the water incorporated? Well, there are two ways:

1. Attach to a metal center.
2. The water crystallizes within the complex.

Let's break this down:

Attach to a metal center

The first way is by attaching to a metal center. Let's use Epsom salt (MgSO₄) as an example. Here is what that structure looks like with no water:

Fig.1-Anhydrous Epsom salt

Now let's take a look at when water is added/absorbed:

Fig.2-MgSO4 incorporates water into its structure

Here, we see that the water has surrounded itself around the metal center. Instead of magnesium (Mg²⁺) and sulfate (SO₄^2-) being bonded, it is the magnesium complex and the sulfate ion.

Water crystallizes within the complex

Sometimes, instead of the water binding to the metal center, it hydrogen bonds to the metal complex. Hydrogen bonding is an intermolecular force of attraction, where the partially negative oxygen is attracted to a species with either a partial or full positive charge and the partially positive hydrogen is attracted to a species with either a partial of full negative charge.

Typically, hydrates that have water crystallized within the complex also have water molecules bound to the metal center, like in FeSO₄ · 7H₂O, shown below:

Fig.3-Iron (II) sulfate heptahydrate

Six of the water molecules (shown as the red ball (oxygen) and the two white balls (hydrogen) are bound to the iron (orange) center. However one the of water molecules is not bound, and instead is attracted to the metal complex and sulfate (yellow ball (sulfur) bonded to oxygens) by hydrogen bonding.

The water molecules that are not bound to the metal center are called water of hydration or water of crystallization

Properties of Hydrates

Most of the time, you can heat the hydrate to get rid of the water of hydration. The structure and texture of the anhydrous compound, which is what's left after heating, will be different from those of the hydrate. It may also be a different color.

For example, the sample below shows anhydrous copper (II) sulfate gaining water to become copper (II) sulfate pentahydrate.

Fig.4-Copper (II) sulfate changes color (white to blue) as it takes in water

Generally, the following are true of any anhydrous compound made from a hydrate:

1. Easy to dissolve in water (highly solvable in water).
2. The color of the anhydrous compound will be similar to the color of the original hydrate when it is dissolved in water, even if the color changed when it went from the hydrate to the anhydrous compound.

At room temperature, most hydrates are stable. However, some hydrates will lose their water when the vapor pressure of a hydrate is greater than the vapor pressure of the air. This process is called efflorescence.

Other substances can take in water from the air on their own (called hygroscopic). Some hygroscopic substances, like P₂O₅ and anhydrous CaCl₂, are often used to "dry" liquids and gases. These substances are called desiccants. This is why you often have silica gel packets in an object's packing to prevent it from taking in water and getting damaged.

Other hygroscopic things, like solid NaOH, take in so much water from the air that they dissolve in water. These things are called deliquescent.

Hydrates Formula

So how do we know how much water is even in these substances? To get the formula of a hydrate, we compare the weights of the hydrate and the anhydrous solid. The mass of water that evaporated can be found by taking the mass of the original hydrate and subtracting the mass of the dry solid:

$$m_{H_2O}=m_{Hydrate}-m_{Anhydrous\,solid}$$

The next step is to divide the mass of water by its molar mass to get the number of moles:

$$n_{H_2O}=\frac{m_{H_2O}}{MM_{H_2O}}$$

$$n_{Anhydrous\,solid}=\frac{m_{Anhydrous\,solid}}{MM_{Anhydrous\,solid}}$$

Lastly, we divide the larger molar amount by the smaller molar amount to get the ratio of anhydrous solid to water.

$$x=\frac{n_{H_2O\,or\,Anhydrous\,solid}}{n_{Anhydrous\,solid\,or\,H_2O}}$$

To better understand this, let's work on an example.

Types of Hydrates

There are three kinds of hydrates:

1. Inorganic.
2. Organic.
3. Gas (or clathrate) hydrates.

Inorganic Hydrates: Inorganic hydrates are the ones we have been discussing so far. In inorganic hydrates, the water molecules are only loosely attached to the compound, and there is no chemical reaction. The water molecule(s) can be easily taken out of the compound, such as by heating it. "Anhydrous" refers to an inorganic hydrate that has lost all of its water molecules. Most hydrates are made up of inorganic hydrates.

Organic Hydrates: Organic hydrates are formed by hydration, which is the addition of water or its components to an organic compound through a chemical reaction. Like with inorganic hydrates, some organic hydrates can be formed without altering the chemistry of the base molecule (i.e., water of hydration)

Gas Hydrates (Clathrate): In gas hydrates, the gas molecule, which is usually methane, is held in place by a loose framework made of water molecules. The "cage" of water molecules is called the host, while the gas inside is called the guest molecule.

Below is an example of a methane hydrate:

Fig.5-A methane gas hydrate

The water molecules(in blue) form a cage with a large void in the center for the "guest" to occupy. In this case, the guest molecule is methane (CH₄, shown in red (hydrogen) and black (carbon)).

Naming System of Hydrates

When writing formulas for and giving names to inorganic hydrates, there are rules to follow. Because the water molecules aren't part of the real structure of the compound, this changes how chemical formulas for inorganic hydrates are written.

Naming hydrates is relatively simple. First, you write the name of the main compound as you usually would. Then, you write a prefix + hydrate. The prefix is based on the number of moles of water present per 1 mole of the solid.

The table below shows the prefixes for different numbers of water molecules.

Some examples of the names of hydrates and their formulas:

• CuSO₄⋅ 5H₂O : Copper(II) sulfate pentahydrate (The Roman numerals in the name show the charge of the metal. Here, the copper ion has a charge of 2+).

• CoCl₂•6H₂O: Cobalt(II) chloride hexahydrate.

• BeSO₄ ⋅ 4H₂O: Beryllium sulfate tetrahydrate.

Commonly used Hydrates

Now that we've covered the basics of hydrates, we are going to take a look at some common hydrates you may see in your everyday life. These are:

1. Epsom salts.

2. Washing soda.

3. Borax.

4. Copper (II) sulfate.

5. Cobalt (II) chloride.

Epsom Salts

Fig.6-Epsom Salts

Formula: MgSO₄ ⋅ 7H₂O

Uses: Epsom salts can be used to soothe sore muscles, as bath salts, to lower systolic blood pressure, and to help plants grow by adding them to the soil.

Washing Soda

Fig.7-Washing Soda

Hydrate Name: Sodium carbonate decahydrate.

Formula: Na₂CO₃⋅10H₂O

Uses: Washing soda was one of the first kinds of soap. Even now, it is sometimes used as a cleaner. Washing soda is efflorescent, which means that at room temperature, it loses some of its water molecules.

Borax

Fig.8-Borax

Hydrate Name: Sodium tetraborate decahydrate.

Formula: Na₂B₄O₇⋅10H₂O

Borax is used in many cleaning products, cosmetics, enamel glazes, and fire retardants.

Copper (II) Sulfate

Fig.9-Copper (II) sulfate

Hydrate Name: Copper (II) sulfate pentahydrate.

Formula: CuSO₄⋅ 5H₂O

Use: When copper sulfate is mixed with water, it turns a bright blue color and has been used as a dye in paintings and pottery. It has also been used as a fungicide and herbicide.

Cobalt Chloride

Fig.10- Cobalt (II) Chloride

Hydrate Name: Cobalt (ll) chloride hexahydrate.

Formula: CoCl₂⋅ 6H₂O

Uses: When it has water, cobalt chloride is purple, but when it doesn't have water, it is light blue. Cobalt chloride-coated papers are sold as a way to find out if something is wet. The papers are blue while they are in the vial, but if moisture is found, they will turn pink once they are taken out.